Energy (E) is the ability to do work.
Many types, but we can say 3 main
US Energy Consumption by Source
Radiant Energy
Light Energy
 Visible
and Invisible
Travels in waves over
 Electromagnetic
 Waves
that spread out in
all directions from the
 Visible light, UV light, Infra
Red Radiation, X-rays,
microwaves, radio waves
Potential Energy (PE)
Stored Energy
 Due
to position
 Chemical
 Nuclear
 Fuels
 Attractions
Kinetic Energy (KE)
Energy of motion
 Atomic
 Molecular
 Vibration
 Rotation
 Translation
 Movement
Kinetic Energy
Can be calculated:
How are each type shown here?
 Rainbow
 Windmill
= visible light
 All
molecules store
 Water
in clouds
 Air
 Materials
the windmill is
made from, the plants
at the bottom
Temperature Scales: Measuring that
Thermal Energy
Fahrenheit (oF)
Celsius (oC)
A Note on the Fahrenheit Scale
NEVER use it in this class. Ever. Only Belize and the US use this
Gabriel Fahrenheit made great thermometers. His scale was
replicated the world over because of this. But if you stop and
think about it, does 32°F for freezing make sense, or 212°F for
boiling? 180 degrees separates them.
100 degrees, as in the Celcius scale (sometimes called the
Centigrade scale) makes much more sense.
Fahrenheit based 0°F on the freezing point of water mixed with
NH4Cl, and 32°F for freezing water, and 96°F for human body
temperature (he was off by 2.6°). Why? Because he felt like it
and it was easy to draw lines at those intervals.
(According to a letter Fahrenheit wrote to his friend Herman Boerhaave,
[8] his scale was built on the work of Ole Rømer, whom he had met earlier. In Rømer’s scale, brine freezes at 0 degrees, ice
melts at 7.5 degrees, body temperature is 22.5, and water boils at 60 degrees. Fahrenheit multiplied each value by four in order to eliminate fractions and increase the granularity of the scale. He
then re-calibrated his scale using the melting point of ice and normal human body temperature (which were at 30 and 90 degrees); he adjusted the scale so that the melting point of ice would be 32
degrees and body temperature 96 degrees, so that 64 intervals would separate the two, allowing him to mark degree lines on his instruments by simply bisecting the interval six times (since 64 is 2 to
the sixth power). I took this from Wikipedia.
Kelvin Temperatures
Based on absolute zero (0 K, -273 oC)
 The
temperature at which ALL KE stops
 NO molecular motion.
 Lowest temperature theoretically possible
 Can’t
really get there in real life
3rd Law of Thermodynamics in a few
 (See
K = oC + 273
 Technically
273.14, but we can stop at 3
significant digits
Why do we need the Kelvin scale?
Two reasons
We need a scale that is relative
to molecular motion
for certain topics
 You
can’t use negative numbers to indicate motion when it
IS present
 -20°C makes NO sense in light of indicating motion
 And 40°C ISN’T twice as much motion as 20°C,
(40K IS twice the motion of 20K)
Because when working with equations, can’t use zero
We get undefined answers if we divide
 We get answers of 0 if we multiple
 And those answers would NOT make sense if compared to answers
calculated with a positive or negative number
The 4 Es:
Energy, Exergy, Entropy, & Enthalpy
Energy (E): The
ability to do work
Entropy (S): The
measure of the
disorder of a
There will be more on these!
Exergy: The energy
available to do
 No symbol
Enthalpy(H): The
thermal energy
(heat) content of a
The study of energy flow
 inter-relation
between heat, work, and energy of a
Summary of the three laws:
The energy in the universe is constant
Things get more disorganized over time in a system
until everything is equal
You can’t reach absolute zero
Law of Thermodynamics
The energy in the universe is constant
 E=mc2
 Law
of Conservation Matter
Matter can not be created or destroyed
 Law
of Conservation of Energy
can not be created or destroyed
 However,
matter and energy can both change
forms in chemical reactions
 Can also interconvert between matter and
energy in NUCLEAR reactions (more on this later
this year.)
Summed up: You can not win. You can’t get something for nothing because
energy and matter are conserved.
Before the
Entropy (S) is a measure of
DISORGANIZATION in a system (this simply
put; there is a much more complicated
description about the unavailable energy to do
 Anything
disorganized has higher entropy than
something organized
Exergy is the Energy available to do work
Law of Thermodynamics
Things get more disorganized over time in a
system until everything equilibrium is reached
(everything is equal)
 Heat
flows from hot to cold, not the reverse
 Law of Entropy
By nature, things get more disorganized to
spread out energy and matter
The quality of the energy (which is exergy)
decreases over time
Summed up: You can not break even. You can not return to the same
energy state because things get more disorganized (gain entropy)
Exergy and Energy
The energy of the universe is constant, but exergy is constantly
consumed. This can be compared with a tooth-paste tube: When you
squeeze the tube (= conduct any process) the paste (= exergy)
comes out. You can never put the paste back in the tube again (try!),
and in the end you have only the tube itself (= low-exergy) left.
When you squeeze the tube, the depressions (= entropy) will
increase. (The entropy of a system increases when exergy is lost) But
you can never take the depressions in the tube and 'un-brush' your
teeth. (I.e. entropy is not negative exergy.)
When you buy energy from the electricity network, you actually buy
exergy. You can find the energy as room temperature heat after
some time, but you can not take that room temperature energy back
to the electricity company and ask for money back. They won't
accept it.
Energy and Matter Gain Entropy Over Time
Exergy: The Energy available to do
Law of Thermodynamics
You can’t reach absolute zero and
expect things to happen
absolute zero, all kinetic motion
ceases. And that energy needs to go
somewhere. It goes to something else.
And gets transferred back until
everything is at an equal temperature.
Summed up: You can not get out of the game, because
absolute zero is unobtainable.
Law of Conservation of Energy
Energy cannot be created or
destroyed…but it CAN change forms.
 Example:
Burning wood in a fire
The energy in chemical bonds is released
as heat (KE and PE), light (RE), sound (KE)
 These
 have
forms of energy are less useful
less exergy
Radiant Energy:
EM Waves
The CPE in these items could:
Rio Summer Olympics
Proposed Solar
Combinations of PE and KE are very
common on a large scale
KE and PE animation
PE and KE
When E changes forms…
The amount of energy one thing loses is
gained somewhere else.
lost = E gained (Law of Conservation of
 But the E gained is usually not all in one
place (2nd Law of thermodynamics)
 It
is spread out (more entropy)
 Often
in the forms of heat and light
 Which are less useful (less exergy)
Energy Transformations
Thermal Energy: KE + PE on the small scale
• What’s up with Temperature vs Heat?
• Temperature is related to the average kinetic energy
of the particles in a substance.
Thermal energy relationships
As temperature increases, so does thermal
energy (because the energy of the
particles increased).
If the temperature stays the same, the
thermal energy in a more massive
substance is higher (because it is a total
measure of energy).
Cup gets cooler while hand
gets warmer
The flow of thermal
energy from one
object to another.
Heat always
flows from
warmer to
cooler objects.
Ice gets warmer
while hand gets
Heat and Temperature
Heat: the measure of the flow of
RANDOM kinetic energy
 Temperature: the measure of heat
 So…temperature is a measure of
kinetic energy of the particles of a
* Sometimes heat is radiated as IR (infra-red radiation,
a form of radiant energy)
Thermal Energy
•Thermal Energy is the
total of all the (kinetic and
potential) heat energy of
all the particles in a
•PE from how the molecules are placed relative to
each other (attractions)
•Farther = more PE, just like how something farther
off the ground has higher gravitational PE
Exothermic and Endothermic Processes
Energy is being
gained/ absorbed by
the object or substance
(called the system)
from the surroundings
Have positive change
in enthalpy values
Energy is lost/
released from the
object or substance
(called the system) to
the surroundings
Have negative change
in enthalpy values
The big picture…
How do we see this energy cycling in the
real world, and not just as a part of
Chemistry class?
 Around
the house?
 In the environment?
 While thinking about a car?
•If the cup is the system, it is
undergoing an exothermic
process because it is losing
heat to the surroundings
•If the ice is the system, it
is undergoing an
endothermic process
because it is absorbing
heat from the surroundings
Cup gets cooler while
hand gets warmer
Ice gets warmer
while hand gets
Which is process is endothermic? Which
is exothermic?
Trophic Levels and Energy
are all
Carnivores and
Carnivores and
1°Consumers: Herbivores
Producers: Autotrophs
Energy Out;
90% per level
Can the world really run out of
World-Wide Energy Sources, (2007)
Phase Diagrams
Tell what state of matter a material is in at a given
temperature and pressure
The triple point is the pressure and temperature when a
solid, liquid, and a gas of the same substance exist at
Equilibrium: When there is no net change
Here referring to changes in state
 Can also refer to temperature and chemicals
The critical point is the temperature above which a
substance will always be a gas, regardless of pressure
Fullerton Phase Diagram Explorer Link
Phase Diagrams
Phase Diagram for Water
A few terms
Freezing Point - The temperature at which the solid
and liquid phases of a substance are in equilibrium
at atmospheric pressure.
 The
same temperature as the melting point
Boiling Point - The temperature at which the vapor
pressure of a liquid is equal to the pressure on the
Vapor Pressure- The pressure at which the
vaporization rates are equal to condensation rates
Phase Changes
The heat
energy) content
of a system
States of Matter and Entropy
The states are NOT plateaus because entropy is NOT constant.
This isn’t a phase change diagram.
Energy and Matter and Connected
Any change in
by a change in
Phase Changes and Energy
Temperature, ̊C
Heating Curve
Why does temperature remains
constant when melting or boiling?
During melting or boiling, energy is
from the surroundings
Due to the increase in the thermal energy of the particles
from the increase in PE of the particles
 Molecules are
moving apart
 breaking attractions which
 Absorbs latent (hidden) heat
 can not be measured on a thermometer
 Substance
(system) gets warmer
The E’s and Heating
•Endothermic process
•Energy is absorbed from surroundings
•Entropy increases
•Enthalpy is positive (+ΔH) since heat added
•Exergy decreases
Why does temperature remains constant
when freezing or condensing?
During freezing or condensing, energy is
released to the surroundings
 Due
to the decrease in the thermal energy from the
decrease in PE of the particles
 Molecules are
moving closer
 forming new attractions that are
 Releasing latent (hidden) heat
can not be measured on a thermometer
 Substance
(system) gets colder
The E’s and Cooling
•Exothermic process
•Energy is lost to surroundings
•Entropy decreases
•Enthalpy is negative (-ΔH) since heat is lost
•Exergy increases
What happens during each segment
Cooling Curve: The Reverse of a
Heating Curve
Temperature, ̊ C
Time, minutes
Measuring the Energy of Phase
The math of thermal energy flow
REMEMBER: Energy and Matter and
Any change in
matter ALWAYS
is accompanied
by a change in
 This includes
changes in
and/ or phase
Specific Heat : c
• Things heat up or cool down at different
Land heats up and cools down faster than water,
and aren’t we lucky for that!?
•Specific heat is the amount of heat required to
raise the temperature of 1 kg of a material by one
degree °C
•cwater = 4.184 J / g °C
•the number is high; water “holds” its heat
•c sand= 0.664 J / g °C
•less E than water to change it; it doesn’t hold
heat as well as water does
This is why land heats up quickly during the day and
cools quickly at night and why water takes longer.
Why does water have such a high
specific heat?
Water molecules form strong attractions with other water molecules;
it takes more heat energy to break those attractions than other
materials with weaker forces of attraction between them.
Specific Heat Capacities of Selected
 cwater
= 4.184 J / g °C
 cice = 2.09 J / g °C
 csteam = 1.99 J / g °C
 csand = 0.664 J / g °C
 cAl = 0.90 J / g °C
 cFe = 0.449 J / g °C
Heat can be Transferred even if there is No
Change in State
q = mc∆T
Remember this? Which is process is
endothermic? Which is exothermic?
Now we care about how much energy
is being transferred, and are ready to
calculate that change.
Calculating Changes in Energy: The
Calorimetry Equation
q = mcT
•q = change in thermal energy
•(+) value means heat is absorbed
•(-) value means heat is released
•m = mass of substance
•T = change in temperature (Tfinal – Tinitial)
•c = specific heat of substance
•Each substance has a different c (see CRH, p__)
•Different states of matter for the same substance may have
a different c
Specific Heat Capacity Problems
If 25.0 g of Al cool from 310 oC to 37 oC, how many joules of heat energy are lost
by the Al?
Notice that the negative sign on q signals heat “lost by” or
transferred OUT of Al.
Was this an endothermic or exothermic process?
Or… Heat Transfer can cause a
Change of State
Changes of state involve energy changes at constant T
Ice + 334 J/g (heat of fusion) -----> Liquid water
Is there an equation? Of course!
Or… Heat Transfer can cause a
Change of State
Changes of state involve energy at constant T
H20(s) +334 J/g  H20(l)
Ice + 334 J/g (heat of fusion)  Liquid water
q = mΔHfusion
•m= mass
•ΔHfusion = the enthalpy of melting
• the change in thermal energy associated
with melting
•Units are J/g or KJ/Kg
q = mΔHfusion
q = mc∆T?
Well, when a phase changes THERE IS NO
change in temperature… but there is definitely
a change in energy!
Sample Problem:
How much heat energy is required to melt 25.0g of ice,
(assuming constant temperature of O°C)?
Value is positive, which means heat is absorbed, which makes
Latent heat* and the PE of particles
strong attraction
breaks up
weak attraction
*Latent means hidden. Latent heat is the thermal energy
(potential energy) associated with the attractions between
molecules, and can not be measured with a thermometer.
Latent heat and the PE or particles
Energy has to be supplied to oppose the
attractive force of the particles.
PE  as
PE related to the forces of attraction between the particles
solid  liquid or liquid  gas
average potential energy 
Latent heat and PE
The transfer of energy does not change the KE.
Temperature does not change.
latent heat = change in PE between
molecules during change of state
Video and song:
Energy changes accompany changes in state; either:
Energy is added (endothermic)
• Gain thermal energy
• Molecules
• Move more (gain KE)
• Separate (gain PE from broken attractions between
• Have a higher entropy
• Are more disorganized
Energy is removed (exothermic)
• Molecules move less
• Lose thermal energy
• Move less (lose KE)
• Move closer (lose PE from new attractions between
• Have lower entropy
• Get more organized
Latent Heats
You have a certain energy change associated with
changing state. These values are usually reported
for fusion and vaporization as:
 ΔHfusion= (latent)
Heat of fusion (melting)
 Δ Hvaporization = (latent) Heat of vaporization
 Δ Hsublimation =(latent) Heat of sublimation
Different materials have different values for each
What about freezing and
Values for freezing and condensation are not
typically listed, but are the negative values of those
for fusion and vaporization because the energy
transferred is the same, but in the opposite direction
 (latent)
Heat of freezing= -ΔHfusion
 (latent) Heat of condensation= -Δ Hvaporization
Enthalpy changes with phase changes
Enthalpy values for H2O
∆Hfusion= 334 J/g
 ∆Hvaporization= 2259 J/g
 ∆Hsublimation = 2594 kJ/g
Summing it all up: How do you know what to do to
calculate energy changes?
Check to see if there is a temperature change.
If yes, use q=mcΔT.
Also, check to see if there is a phase change.
If yes, you need to use
q= Δ Hfusionmass
• or
q= Δ Hvaporizationmass
depending on which one applies*
or both if there are two phase changes
*If the material freezes or condenses. You can use the negative
value Δ Hfusion or Δ Hvaporization
How much energy is required to change 0.5 kg of
water at 0 °C to ice?
Things you know:
•m =
•There is____ temperature change, and there is
change of state (freezing)
•The water is going __________
this all tells you to use _________(negative of
melting value) in
(The negative value makes sense since you are
cooling the water, so energy leaves)
How much energy is required to melt 0.5 kg of ice at 0 °C
temperature raised to 80 °C?
Total energy required
Heat & Changes of State
What quantity of heat is required to melt 500. g of ice and
heat the water to steam at 100. oC?
Heat of fusion of ice = 334 J/g
Specific heat of water = 4.184 J/g•°C
Heat of vaporization = 2259 J/g
+334 J/g
Putting it all together…
So… if I want the total heat to take ice and turn it to steam I need to add values from
3 steps…
1. To melt the ice I need to multiply the heat of fusion with the mass
• q = ∆Hfusionm
Then, there is moving the temperature from 0°C to 100°C.
• For this there is a change in temperature so we use
q= mc∆T
That just takes us to 100°C, what about vaporizing the molecules?
We need q=∆Hvaporizationm
Add up all the values, and you have it.
(However, if you are taking it from below the freezing point to above 100°C, you
need to add in the changes with q=mc ∆ T there, too!)
And now… More! Heat & Changes of State
How much heat is required to melt 500. g of ice and heat the water to
steam at 100 oC?
To melt ice
2. To raise water from 0 oC to 100 oC :
3. To evaporate water at 100 oC:
4. Total heat energy =
Maybe a picture can help….
Putting it all together:
How are matter and energy related?
What influences does energy have on
matter? What does this tell us about the
world as we know it?
Making Pizza: Changing Matter
Describe the pizza making process in terms of:
 Matter
 States
(s, l, g)
 Elements, compounds, mixtures
 Homogeneous
and heterogeneous mixtures
Properties and changes
 Both
 chemical
and physical
 Intrinsic (intensive) and extrinsic (extensive)

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